The History and Use of Our Earth's Chemical Elements: A Reference Guide, Second Edition

Our Earth's Chemical Elements: A Reference Guide, Second Edition

Robert E. Krebs

Greenwood Press

The฀History฀and฀฀ Use฀of฀Our฀Earth’s฀ Chemical฀Elements

A Reference Guide Second Edition

Robert฀E.฀Krebs Illustrations฀by฀Rae฀Déjur

GREENWOOD PRESS

Westport, Connecticut • London

Library of Congress Cataloging-in-Publication Data

Krebs, Robert E., 1922 The history and use of our earth’s chemical elements : a reference guide / Robert E. Krebs ; illustrations by Rae Déjur. — 2nd ed.

p. cm. Includes bibliographical references and index. ISBN 0–313–33438–2 (alk. paper)

1. Chemical elements. I. Title. QD466.K69 2006 546—dc22 2006012032

British Library Cataloguing in Publication Data is available. Copyright © 2006 by Robert E. Krebs All rights reserved. No portion of this book may be

reproduced, by any process or technique, without the express written consent of the publisher.

Library of Congress Catalog Card Number: 2006012032 ISBN: 0–313–33438–2 First published in 2006 Greenwood Press, 88 Post Road West, Westport, CT 06881

An imprint of Greenwood Publishing Group, Inc. www.greenwood.com

Printed in the United States of America

∞ TM

The paper used in this book complies with the Permanent Paper Standard issued by the National Information Standards Organization (Z39.48–1984).

To฀Carolyn,฀my฀wife,฀proofreader,฀pre-editor,฀ constructive฀critic,฀and฀friend.

Contents

Alphabetical List of the Elements xv How to Use This Book

ix Introduction

xxiii One

1 The Beginnings of Science

A Short History of Chemistry

1 Origins of the Earth’s Chemical Elements

1 Early Uses of Chemistry

2 The Age of Alchemy

4 The Age of Modern Chemistry

4 Elements and Compounds

5 Two

Atomic Structure

9 Early Ideas of Atomic Structure

9 Some Theoretical Atomic Models

13 The Nucleus and Radiation

16 Chemical Bonding

18 Organic Chemistry

20 Fullerenes and Nanotechnology

25 History

25 Rules for Cataloging the Elements

Guide to the Elements

35 Introduction

Alkali Metals: Periods 1 to 7, Group 1 (IA) 39

Introduction 39 Hydrogen

Alkali Earth Metals: Periods 2 to 7, Group 2 (IIA) 65

Introduction 65 Beryllium

Transition Elements: Metals to Nonmetals 85

Introduction 85 Transition Elements: First Series—Period 4, Groups 3 to 12

87 Scandium

87 Titanium

Contents | ix

Vanadium 93 Chromium

114 Transition Elements: Second Series—Period 5, Groups 3 to 12

143 Transition Elements: Third Series—Period 6, Groups 4 to 12

147 Hafnium

147 Tantalum

150 Tungsten

153 Rhenium

155 Osmium

157 Iridium

159 Platinum

162 Gold

165 Mercury

Metallics—Metalloids—Semiconductors—Nonmetals 173

Introduction 173 The Boron Group (Metallics to Semimetals):

Periods 2 to 6, Group 13 (IIIA) 175 Introduction

186 The Carbon Group (Metalloids to Semiconductors):

Periods 2 to 6, Group 14 (IVA) 189 Introduction

203 The Nitrogen Group (Metalloids to Nonmetals):

Periods 2 to 6, Group 15 (VA) 207 Introduction

220 The Oxygen Group (Oxidizers and Nonmetals):

Periods 2 to 6, Group 16 (VIA) 223 Introduction

223 Oxygen

223 Ozone (Also Group 16)

Contents | xi

Sulfur 234 Selenium

237 Tellurium

239 Polonium

241 The Halogen Group (Nonmetal Oxidizers):

Periods 2 to 6, Group 17 (VIIA) 245 Introduction

257 The Noble Gases (Inert Gas Elements):

Periods 2 to 6, Group 18 (VIIIA) 260 Introduction

Lanthanide Series (Rare-Earth Elements): Period 6 275

Introduction 275 Lanthanum

277 Cerium

279 Praseodymium

281 Neodymium

283 Promethium

285 Samarium

287 Europium

289 Gadolinium

Terbium 292 Dysprosium

Actinide Series (Period 7) and Transuranic Elements * 305

Introduction 305 Actinium

Transactinide Series: Period 7 (Continuation of Actinide Series) 339

Introduction 339 Rutherfordium (Unnilquadium)

342 Dubnium (Unnilpentium)

343 * The following elements belong to the subseries of transuranic elements as well as to the actinide series: Neptunium,

Plutonium, Americium, Curium, Berkelium, Californium, Einsteinium, Fermium, Mendelevium, Nobelium, Lawrencium.

Contents | xiii

Seaborgium (Unnilhexium) 345 Bohrium (Unnilseptium)

346 Hassium (Unniloctium)

349 Meitnerium (Unnilennium)

350 Darmstadtium (Ununnilium)

352 Röentgenium (Unununium)

352 Ununbium ( 112 Uub-285)

353 Ununtrium ( 113 Uut-284)

354

The Superactinides (Super Heavy Elements) and Possible Future Elements 357

Introduction 357 Ununquadium ( 114 Uuq)

358 Ununpentium ( 115 Uup)

359 Ununhexium ( 116 Uuh)

361 Ununseptium ( 117 Uus)

362 Ununoctium ( 118 Uuo)

363 Possible Future Elements

365 Glossary of Technical Terms 367

Appendix A: Discoverers of the Chemical Elements 395 Appendix B: Chemical Elements in Order of Abundance

in Earth’s Crust 403 Appendix C: Nobel Laureates in Chemistry (1901–2005) 407 Selected Bibliography 413 Index 417

Alphabetical List of the Elements

Chemical

Atomic

Name of Element

Symbol

Number Atomic Weight Page

Be 4 9.012182 65 Bismuth

C 6 12.01115 189 Cerium

279 Cesium

Ce 58 140.116

55 132.90546 59 Chlorine

Cs

248 Chromium

Cl

95 Cobalt

Cr

105 Copper

Co

Cu

Chemical

Atomic

Name of Element

Symbol

Number Atomic Weight Page

F 9 18.9984 245 Francium

Gd 64 157.25 290 Gallium

Ga 31 69.723 181 Germanium

Ge 32 72.64 198 Gold

79 196.967 165 Hafnium

Au

Hf 72 178.49 147 Helium

He 2 4.002602 261 Holmium

67 164.903 295 Hydrogen

Ho

H 1 1.0079 40 Indium

49 114.818 159 Iodine

In

I 53 126.9044 254 Iridium

77 192.217 159 Iron

Ir

Fe 26 55.847 100 Krypton

Hg 80 200.59 168 Molybdenum

42 95.94 127 Neodymium

Mo

60 144.24 283 Neon

Nd

10 20.179 265 Neptunium

Ne

93 237.0482 316 Nickel

Np

28 58.6934 108 Niobium

Ni

41 92.906 124 Nitrogen

Nb

7 14.0067 207 Nobelium

No

Alphabetical List of the Elements | xvii

Chemical

Atomic

Name of Element

Symbol

Number Atomic Weight

Page

Osmium

157 Oxygen

Os

223 Palladium

137 Phosphorus

Pd

212 Platinum

162 Plutonium

Pt

318 Polonium

Pu

241 Potassium

Po

53 Praseodymium

281 Promethium

Pr

285 Protactinium

Pm

311 Radium

Pa

81 Radon

Ra

272 Rhenium

Rn

155 Rhodium

Re

135 Rubidium

Rh

57 Ruthenium

Rb

133 Samarium

Ru

287 Scandium

Sm

87 Selenium

Sc

34 78.96 237 Silicon

Se

194 Silver

Si

140 Sodium

Ag 47 107.868

50 Strontium

Na

38 87.62 76 Sulfur

Sr

234 Tantalum

150 Technetium

Ta

130 Tellurium

Tc

239 Terbium

Te

292 Thallium

Tb

186 Thorium

Tl

309 Thulium

Th

299 Tin

Tm

200 Titanium

Sn

22 47.88 90 Tungsten

Ti

153 Uranium

312 Vanadium

93 Xenon

V 23 50.9415

Xe 54 131.293

Chemical

Atomic

Name of Element

Symbol

Number Atomic Weight Page

˜253 to 263 342 Dubnium

Rf (Unq)

Db ( Unp)

Sg (Unh)

Bh (Uns)

Hs (Uno)

Mt (Une)

350 Röentgenium

Ds (Uun)

Rg (Uuu)

˜292 (in dispute) 361 Ununseptium

Uuh

Unconfirmed (undiscovered) 362 Ununoctium

Unconfirmed (in dispute) 363

Note: Superactinides and super heavy elements (SHE) are elements beyond lawrencium 103. All are artificially produced, unstable, and radioactive and have very short half-lives. Most are made in small amounts, even one atom at a time.

How to Use This Book

The History and Use of Our Earth’s Chemical Elements is organized to reflect the chemical and physical properties of the elements that are depicted in the periodic table of chemical elements. This book uses the same general format as the periodic table. Familiarity with how the periodic table is organized and its terminology is not required to understand this book. However, the beauty of the table’s organization will become apparent as this book assists you in understanding how the elements are classified according to their atomic numbers, atomic weights, and other chemical and physical properties.

There are several ways to use this book. First, the book is not divided into chapters, but rather into sections. The first five sections include background information that will provide an understanding of basic chemistry and how the rest of the book is organized. If you already are familiar with basic chemistry, you may wish to skip some of the beginning sections. There are three ways to proceed if you are just interested in knowing more about a particular ele- ment.

1. One way is to look up the name of the element in which you are interested in the “Alphabetical List of the Elements” located immediately following the table of contents. 2. Another way to start is to look up the element you wish to know more about in its period or group as listed in the table of contents. The Contents is really the guide for presenting the major classes or categories of elements, that is., periods and groups of elements.

3. Still another way is to look up the page for the element in which you are interested in the index at the end of the book.

The first sections of this reference book set the stage for the presentation of the elements. First is the section “How to Use This Book” followed by a short introduction. Next is “A Short History of Chemistry,” the narrative of which progresses from prehistoric times to the Age of Alchemy and then to the Age of Modern Chemistry. Next is the section titled “Atomic Structure,” which traces the history of our knowledge of the structure of the atom; some theo- retical models, including quantum mechanics; the discovery of subatomic (nuclear) particles The first sections of this reference book set the stage for the presentation of the elements. First is the section “How to Use This Book” followed by a short introduction. Next is “A Short History of Chemistry,” the narrative of which progresses from prehistoric times to the Age of Alchemy and then to the Age of Modern Chemistry. Next is the section titled “Atomic Structure,” which traces the history of our knowledge of the structure of the atom; some theo- retical models, including quantum mechanics; the discovery of subatomic (nuclear) particles

Periodic Table of the Chemical Elements,” which describes the table’s organization and how to use the table. The rules for placing the elements within the matrix structure of the periodic table are given in this section.

“A Guide to the Elements,” which is the major section of the book, presents the elements as they are arranged in the periodic table. Each element is presented as a separate entry that includes information on

• the element’s symbol, period, and group and where it is found in the table, its atomic num- ber, its atomic mass (weight), its valence and oxidation state, its natural state, the origin of its name, and the element’s isotopes;

• a figure representing the electron configuration for each element up to and including ele- ment 103 (lawrencium); • the important chemical and physical properties of each element; • the important characteristics of each element; • each element’s abundance and its source on Earth; • the history of each element, including the discoverer(s) of the element and how and when

it was discovered; • a section for each element’s important uses; • examples of common compounds; and • an explanation of each element’s potential hazards to humans and the environment.

As previously stated, the reader can locate specific elements by checking the Contents for the elements’ placement in the periodic table or by consulting the “Alphabetical List of the Elements” that follows the Contents.

The following notations are used in this book: BCE = Before the Common Era (instead of B.C.)

CE = Common Era (instead of A.D.) BP= Before the Present (time)

c. = Approximate date (e.g., approximate birth or death dates) ˜ = Approximate amount, quantity, or figure amu = atomic mass units, or for heavy elements average mass units. aw = atomic weight mw = molecular weight

0 n-1 = a single neutron (no charge with atomic weight of 1)

α฀= alpha particle (helium nucleus; 2 He-2 or He ++ )

β฀= beta particle (high-energy electron) λ฀= gamma radiation (similar to high-energy X-rays) SF = spontaneous fission SHE = abbreviation for super heavy elements

Z = atomic number ( 92 uranium, or 92 U, can be shown as Z-92)

IUPAC = International Union of Pure and Applied Chemistry (England)

How to Use This Book | xxi

ACS = American Chemical Society (U.S.A.) SHIP = Separator for Heavy Ion Products (Germany) JINR = Joint Institute for Nuclear Research (Russia) ppm = parts per million ppb = parts per billion % = percent ∆= heat ↑ = gas produced → = yields

Terms set in bold type throughout the text can be found in the Glossary of Technical Terms.

Introduction

This volume is a reference for students and other readers interested in chemistry as well as for school and public library use.

It provides the background of how we came to know and understand the chemical nature of our planet and everything in the universe, including ourselves. Early humans had limited knowledge and use of their chemical environment. Through the ages humans progressed from

a practical to a spiritual and finally to a rational approach to the nature of the chemical ele- ments found on Earth, as well as those in the entire universe. As science developed, our accumulation of knowledge about the structure of atoms and molecules was an achievement of early philosophers and scientists. These men and women did not use scientific procedures, but they did build the foundation of our current understanding of the structure of matter and how different species of matter interact. This history has led to our current understanding of the theoretical and practical nature of the chemical elements.

This reference book describes the chemical elements according to such characteristics as structure, size, weight, activity (energy), abundance, usefulness, and hazards. Each element’s structure relates to its “fit” within the periodic table of the chemical elements. The book is about chemical elements found on Earth as well as in the entire universe. Interestingly, the proportions (by weight or number) of elements found on Earth are not the ratios found in the rest of the universe. For instance, although hydrogen is the most common element in the universe (90+%), the most common element on Earth is oxygen (49%), which was formed by phototropic bacteria and later green plants long after the Earth was formed about 4.5 bil- lion years ago.

Chemistry is a physical science that studies the structure and properties of elementary mat- ter. Matter interacts with other substances. Matter can be defined as something that occupies space, has mass, and cannot be created or destroyed, but can be changed from one form to another. Matter is “stuff” that can be perceived by one or more of our senses, as opposed to something intangible such as an idea, the mind, or spirits (ghosts and angels). We usually think of matter as the chemical elements composed of atoms and the compounds composed of molecules, which are combinations of atoms. Chemistry is the science of how and what Chemistry is a physical science that studies the structure and properties of elementary mat- ter. Matter interacts with other substances. Matter can be defined as something that occupies space, has mass, and cannot be created or destroyed, but can be changed from one form to another. Matter is “stuff” that can be perceived by one or more of our senses, as opposed to something intangible such as an idea, the mind, or spirits (ghosts and angels). We usually think of matter as the chemical elements composed of atoms and the compounds composed of molecules, which are combinations of atoms. Chemistry is the science of how and what

physical things we can see, touch, and taste. Although chemistry is universal, it is not spiritual or mysterious. We study the chemical and physical nature of what makes up atoms and the interactions of atoms and molecules in the stars, the Earth’s matter, living cells, and everything that exists. Chemistry is also a science of energy. The formation of molecules by the combina- tion of atoms involves energy. Chemical reactions involve energy during the combination or separation of atoms and molecules. When energy is released, the reaction is known as an exothermic chemical reaction. Other chemical reactions that require an input of energy to complete the reaction are known as endothermic reactions. In addition to the nuclei of atoms that are mainly composed of protons and neutrons, there are electrons orbiting the nuclei of atoms. These electrons exist in a specific state, or level of energy. Atomic structure is discussed in the section on atomic structure.

Almost everything we live with on Earth is made up of about 100 different chemical elements that, in various combinations, form compounds. But only a few of these elements are essential to explain life on Earth. The air we breathe, the food we eat, the clothing we wear, the cell phones we use, and our bodies all consist of chemicals. The trillions upon trillions of atoms and molecules in our bodies consist of only six major elements: carbon, oxygen, hydrogen, nitrogen, phosphorus, and sulfur. In addition, there are over 40 trace elements in our bodies that are important for our well-being. Some of these are copper, iron, chromium, magnesium, calcium, and zinc. The current market price for all the elements in your body is about $10.95.

What makes chemistry so interesting is that each specific chemical element is related to its own kind of atom. Elements with specific characteristics have unique atoms. Each type of atom is unique to that element. If you change the basic structure of an atom, you change the structure and properties of the element related to that atom. Also of interest is what happens when two or more different atoms combine to form a molecule of a new substance. Once they form a mol- ecule of a new compound, the original atoms no longer exhibit their original properties.

All this is dependent on the electron arrangement in the shells * around the nucleus of the atom. How atoms interact—that is, how they combine to form molecules—is dependent on the arrangement of their electrons. The electron is one of the three major subatomic particles. It has a negative charge and a negligible mass (only 1/1837 the weight of a proton). Although orbital electrons are continually circling the central positively charged nucleus of the atom, it is not possible to determine exactly where they are within their orbit (shell) at any point in time. (See quantum mechanics in the section on atomic structure) The outer ring, shell, or orbit of electrons consists of a specific number of electrons for each element. For most ele- ments, the electrons in the outermost shell may be thought of as valence electrons because they partially determine the chemical properties of atoms and how they combine with each other. Atoms of some elements may have more than one valence number. Valence is a whole number that represents the combining power of one atom to another and provides the relative amounts of each of the interacting elements in the new molecule of a new compound. Valence is discussed in detail in the section on atomic structure.

* The illustrations that depict the electron configurations of the atoms of each element are based on the Bohr model of quantum energy shells.

Introduction | xxv This reference work uses the periodic table of the elements as the basis for organizing the

presentations of the elements. Once you learn how this remarkable chart is organized, you will be able to relate the characteristics of many elements to each other based on the structure, which can be determined by their placement on the periodic table.

The section “How to Use This Book” provides details and specifics concerning the organi- zation of this user-friendly reference work. Robert E. Krebs

one

A Short History of Chemistry

The Beginnings of Science

There are two major theories of how science developed over the ages. One states that early humans, being curious and having some intelligence, began to explore nature by using trial and error. To continue existing, humans learned what to eat, how to protect themselves, and, when time permitted, how to cope with their environment to make life easier and more understandable. This is the “continuum” or “accumulative” approach to science and discov- ery—which is still ongoing.

The other theory, as presented by Alan Cromer in Uncommon Sense—The Heretical Nature of Science (1993), postulates that science was not a natural sequence or continuum of inventions and discoveries from ancient to modern times. He states that science and technological develop- ments occurred in “spurts” of periods of discovery interspersed with periods of ignorance and status quo. He also states that what we think of as science developed in early Greece, possibly because of the democratic nature of Greece’s culture, which included rational inquiry and debate. From Greece, science spread to Egypt, China, India, and the Mesopotamian region, where it flourished for a time. As Muslims conquered many countries, the Arab world introduced science to southern Europe as far west as Spain. In the Middle Ages, these Arabic texts of Greek science were retranslated first into Latin and later into the vernacular languages of western Europe.

The libraries of Alexandria (in Egypt) became depositories of knowledge and were major contributors to the advancement of science and intellectual studies in many other countries. Modern science is very different from the descriptions of early systems of thought. Early philosophers and theorists lacked the objective methodologies and rational investigative pro- cesses required for the controlled experiments that led to modern science. They were more concerned with seeking universal cures for sickness, transmutation of base metals into gold, and mysticism in general. Most, but not all, ancient philosophers depended more on the writ- ten words of “experts” than on their own observations and insights.

Origins of the Earth’s Chemical Elements

2 | The History and Use of Our Earth’s Chemical Elements from the time they first became cognizant of their environment. Even though it is impossible

to have complete certainty, during the twentieth century and continuing into the present, science has accumulated a tremendous amount of data and knowledge to support the belief that a cataclysmic event now known as the “Big Bang” resulted in the conditions necessary for the creation of the universe and our solar system.

The most accepted theory of several proposed ideas for the origin of the universe, the Big Bang theory states that it all started 13 to 15 billion years ago with the “explosion” of an incredibly small, dense, and compact point source of matter and energy. This resulted in the creation and rapid expansion of space and the beginning of time, as well as the formation of the universe and all that is contained within in it. Astronomers’ observations indicate that the universe continues to expand at an accelerated rate, and this original unknown primordial mass is the source of all the chemical elements and energy existing in the universe. The forma- tion of the chemical elements began in just seconds. Hydrogen was the first element to be formed and remains the most abundant element in the universe. In the heat of the explosion, hydrogen nuclei, having one proton, fused to form helium, which has two protons in its nucleus. Together, hydrogen and helium make up 98 to 99% of all the atoms in the universe. As these gaseous elements condensed, they formed billions of galaxies consisting of trillions of stars as well as dark matter that we can’t see but that constitutes a large portion of the universe. In addition, many of the stars in the galaxies have accompanying planets, comets, meteorites, and cosmic dust as well as all the chemical elements.

Another popular theory speculates that there was no beginning and there will be no end to the universe—it is infinite. Another theory is that there is continuous death and rebirth of stars and matter in the universe. Others postulate that it all started by some spontaneous and unknown force—possibly supernatural.

We know a great deal about the nature of the universe. For instance, the element hydrogen makes up about 75% of all the mass in the universe. In terms of number, about 90% of all atoms in the universe are hydrogen atoms, and most of the rest of the atoms in the universe are helium. All the other heavier elements make up just one to two percent of the total. Interestingly, the

most abundant element on Earth (in number of atoms) is oxygen (O 2 ). Oxygen accounts for about 50% of all the elements found in the Earth’s crust, and silicon, the second most abun- dant element, makes up about 25%. Silicon dioxide (SiO 2 ) accounts for about 87% of the total Earth’s mass. Silicon dioxide is the main chemical compound found in sand and rocks.

Early Uses of Chemistry

Early beliefs about and uses for the chemical elements are not well recorded. Obviously, early humans learned, probably through trial and error, how to use several of the Earth’s chem- ical elements for survival. Fire is one of the earliest examples of the use of a chemical reaction. Humans were obviously aware that fire gave off light and heat and caused death, but in time they learned how to use fire to their advantage, most probably for cooking and warmth. Oil lamps that used liquid fat were invented in about 70,000 BCE. Solid fat to form candles for light originated about 3000 BCE. The burning of a candle involved a great deal of chemistry, but many centuries passed before humans understood the science involved.

An interesting application of early chemistry was the use of fire to make pottery. Early pots were formed from soft clay that, of itself, cannot hold much weight or water. Early humans

3 had baskets, wood containers, and leather bladders—yet there was a need for something in

A Short History of Chemistry |

which to hold water and to cook food. Around 13,000 BCE, people learned, either by accident or by trial and error, to bake the soft clay over fire, causing it to harden. Once fired, the clay pot became a waterproof ceramic, resulting in a more useful container. This may have been the first time humans used fire for a purpose other than for cooking, heat, and protection.

Another example of an early use of chemistry is the discovery of fermentation: overripe fruit or honey, if left uncovered, turned into wine after coming into contact with an airborne organism. During fermentation, large glucose molecules (sugar) break down into the smaller molecules of carbon dioxide and alcohol. The fermentation process that produces alcohol dates back nearly 10,000 years. Wine-making dates back about 7,500 years, evidenced by the residue at the bottom of a clay pot, dated to 5400 BCE, recently found in the mountains of Iran. Wine-making is an ancient chemical process.

A related, and possibly accidental, discovery was bread-making. Presumably some ancient human noticed that when flour pounded out of wild grains was moistened and then exposed to air, it “rose,” forming leavened bread. Airborne yeasts are responsible for the fermentation

that produces the gas carbon dioxide (CO 2 ) in the dough. As the dough rises and then is baked, the carbon dioxide gas produces an open texture in the final product that is far more edible. A standard practice was to save a small portion of the unbaked dough, which contained some of the yeast, as a “starter” for future baking. It was kept separate and cool until mixed with fresh dough for the next batch of bread. Then a new starter was put aside for future bak- ing, and so on. If you heat or bake a starter sample, you run the risk of killing the yeast, which is exactly what happens during the baking process after the bread loaf has been raised by the action of the carbon dioxide and heat.

There are numerous examples of the early uses of chemistry in metalworking. Copper and some other metals were found as nuggets or exposed native deposits. Sometime around 8000 BCE, it was discovered, probably by accident, that when heated, copper ore combines with air

to form a gas (CO 2 ) and metallic copper (Cu). Thus, humans did not need to rely on stone and wood for tools and weapons, given that metal was superior for these purposes and was now available. A later discovery made circa 3600 BCE involved the mixing of copper and tin ores, resulting in the alloy called bronze. This was the beginning of the Bronze Age. Bronze is stronger and holds a sharper edge on tools and weapons than pure copper.

High-grade iron exists in meteorites and in some iron ores. Early humans found these sources of iron but were unable to make use of them except as ornaments. Most iron on Earth exists in ores where it is combined with other substances (often oxygen). It cannot be melted down by wood fires because wood fires do not produce temperatures hot enough to separate the iron from its impurities. In about 1500 BCE, humans learned how to convert wood to charcoal (another chemical reaction) and found that charcoal produces a higher temperature than wood when burned, thus making it possible to smelt the iron from its ore. Iron made sharper, stronger, and more durable tools and weapons than did bronze, thus the beginning of the Iron Age.

These examples, as well as other early uses of chemistry, all involved chemical changes that are well known today. It was many years before humans began studying how chemical changes occur and how to explain and control these reactions.

In 340 BCE, Aristotle (384–322 BCE) published Meteorologica, in which he postulated that the Earth’s matter is composed of four elements—earth, water, air, and fire. His speculations

4 | The History and Use of Our Earth’s Chemical Elements led to the idea that the Earth is composed of “shells,” which is a rather modern concept in

earth science. This was the extent of humankind’s understanding of the composition of the Earth’s chemical elements for several centuries—until the Age of Alchemy.

The Age of Alchemy

Various forms of alchemy were practiced from about 500 BCE into the seventeenth century. It was not until about 320 BCE, at the time of Alexander the Great, that alchemists made serious studies of chemical changes. Alchemy encompassed Greek and Egyptian as well as Arabian and Chinese concepts of matter and energy. The early alchemists were not scientists, and they also did not use scientific procedures, as we think of these terms today. Mostly they were philosophical and theological about their trial-and-error procedures. They were unable to decipher the nature of chemical reactions, but they did make some discoveries that advanced knowledge. In the year 300, Zosimus of Egypt made the first attempt to summarize the accu- mulated knowledge of alchemists.

Unfortunately, most early alchemists are unknown, considering that they were very secre- tive about their methods and left little in the way of written history. Their goals were mystical, economic, secret, unpublished, and unshared. Alchemic practices were also related to medi- cine as well as religion during some periods of time and in some countries. The alchemists’ main search was for the “philosopher’s stone” that could unlock the secrets of transmuta- tion—that is, the secrets of how to transform base metals and chemicals into different, more useful and valuable products, such as gold and silver. This also led to the futile search over many centuries for the elixir vitae that would be both the universal “cure” for all illnesses and the way to achieve immortality.

For over 2,000 years, alchemy was the only “chemistry” studied. Alchemy was the prede- cessor of modern chemistry and contributed to the slow growth of what we know about the Earth’s chemical elements. For example, the alchemists’ interest in a common treatment for all diseases led to the scientific basis for the art of modern medicine. In particular, the alchemist/ physician Paracelsus (1493–1541) introduced a new era of medicine known as iatrochemistry, which is chemistry applied to medicine. In addition, alchemists’ elementary understanding of how different substances react with each other led to the concepts of atoms and their interac- tions to form compounds.

The Age of Modern Chemistry

It is both difficult to determine an exact date for the beginning of modern chemistry and impracticable to bestow the designation of “father of chemistry” on any one individual. Some historians date the end of alchemy and the beginning of modern chemistry to the early seventeenth century. Over the years many men and women of different races and from many countries have contributed to our current knowledge and understanding of chemistry. A few examples follow.

In 1661 Robert Boyle (1627–1691), an early chemist from Great Britain, published a book titled The Skeptical Chymist, which was the beginning of the end of alchemy. His book ruled the perceptions and behavior of early scientists for almost 100 years. Two of his contributions were the use of experimental procedures to determine properties of the chemical elements

5 and the concept that an element is a substance that cannot be changed into something sim-

A Short History of Chemistry |

pler. Robert Boyle is best known for Boyle’s Law, which states that the volume of a gas varies inversely with the pressure applied to the gas when the temperature remains constant. In other words, as you squeeze a container of gas, as occurs when the piston compresses the air and gas mixture in an internal combustion engine, the volume of the gas decreases. In the reverse, if you increase the amount of gas in a closed container, the pressure on the inside of the contain- er becomes greater (similar to pumping up a bicycle tire). On the other hand, if you increase the volume of the container but not the amount of gas, the pressure inside the container will

be reduced. In the early 1700s, Georg Ernst Stahl (1660–1743), a German chemist, devel- oped a theory that when something burned, phlogiston (from the Greek “to set on fire”) was involved. His idea was that burnable things had a limited amount of phlogiston and that when burned they lost their phlogiston, leaving residues that would not burn because they no longer contained this so-called substance. Although this theory was considered viable for many years, it could not be sustained as the physical science of chemistry advanced with new concepts for combustion. Through experimentation, it was shown that different products resulted from combustion, depending on the particular substances that burned in the atmosphere.

Most historians credit the French chemist Antoine-Laurent Lavoisier (1743–1794) with the death of alchemy and the birth of modern chemistry. His many contributions to the profession included the important concept that one must make observational measurements and keep accurate written records. Lavoisier mixed substances, burned common materials, and weighed and measured the results. His work led to the discovery of more than 30 elements. He described acids, bases, and salts as well as many organic compounds. Through a unique

experiment with water (H 2 O), he determined that it is made up of the gases hydrogen and oxygen, with oxygen having a weight eight times that of hydrogen. This led to a later theory of the Law of Definite Proportions, which states that a definite weight of one element always combines with a definite weight of the other(s) in a compound. (It should be noted that Lavoisier was unaware that two atoms of hydrogen combine with one atom of oxygen to form

a water molecule. Therefore, the actual ratio of weight [atomic mass] is 1:16 instead of 1:8.) Up until this time, no standard nomenclature (names and symbols) was used for the elements, compounds, and reactions. In 1769, Lavoisier and others published a book titled The Methods of Chemical Nomenclature that proposed a logical systematic language of chemistry. Even with modifications by the Geneva System of 1892 and additional reforms by the International Union of Pure and Applied Chemistry (IUPAC) in 1930, Lavoisier’s nomenclature of chemi- cal names and symbols is still in use today.

Jöns Jakob Berzelius (1779–1848), a Swedish chemist, is also considered one of the found- ers of modern chemistry. He prepared, purified, and identified more than 2,000 chemical ele- ments and compounds. He also determined the atomic weight (mass) of several elements and replaced pictures of elements with symbols and numbers, which is the basis of our chemical notations today.

Elements and Compounds

Essentially all of us have used the terms “weight” and “mass,” and most people use these terms interchangeably. However, there is an important scientific distinction. The mass of an object is the amount of matter (stuff) contained in a particular body or volume of a substance,

6 | The History and Use of Our Earth’s Chemical Elements regardless of the object’s location in the universe. The mass of an object is constant and always the

same, no matter the planet or galaxy in which it is located.

The weight of an object relates to its size (mass) and distance from the gravitational pull of another body, such as Earth or any other large mass in the universe. In other words, weight is gravity’s effect on the mass of an object. Thus object A’s weight depends on two factors: first, the size (mass) of the two bodies (for instance object A and the Earth); and second, the square of the distance separating the two bodies. Mass is constant in the universe, whereas weight may be thought of as the strength with which gravity “pulls” objects (A) to Earth, or any other object in the universe whose gravity might affect “A.”

The concept of an “atom” is very old. Leucippus of Miletus in Greece (c. 490–430 BCE) proposed that all matter is composed of very minute particles called atoms, from the Greek word atomos, meaning indivisible. Atoms are so small that nothing can be smaller, and they cannot be further divided and still be characteristic of the elements they form. One of Leucippus’ students, the philosopher Democritus (c. 469–370 BCE), is credited with further developing his teacher’s concept of the atom. Another philosopher, Zeno of Elea (c. 495–430 BCE), used paradoxes to present his hypotheses that distance and motion, as well as matter, could be divided into smaller and smaller units in perpetuity. His most famous paradox is the fable of the race between the hare and tortoise wherein the tortoise is given a head start to the halfway point of the total distance of the race before the hare has started. Next, the tortoise progresses to half of the remaining distance—1/4—and then to 1/8, 1/16, and so on, so that no matter how the race progressed, the tortoise would always lead the hare. Democritus chal- lenged Zeno’s concept with the analogy that if a handful of dirt is divided by half, and then that half is divided into half ad infinitum, there would be a final limit to the point of divis- ibility, and this would be the indivisible atom of the dirt (matter). Other Greek philosophers and scientists, including Aristotle and Epicurus (c. 341–270 BCE), accepted Leucippus’ theory of the atom and described it as logical and rational and claimed that it could be used to explain reality and eliminate superstition. Epicurus also stated that atoms were always in constant motion, as well as perceivable, and thus deterministic. Although these ancient scientists lacked experimental evidence, they did speculate philosophically about the concept of atoms, which was not well accepted until much later. It would be many centuries before Epicurus’ idea was developed into the concept of kinetic energy, heat, and thermodynamics.

As mentioned previously, many people contributed to the development of modern chem- istry. An important advancement in our understanding of the Greek philosophers’ concept of the indivisible atom occurred in the early nineteenth century. After conducting many experi- ments, the English chemist John Dalton (1766–1844) published his book New Systems of Chemical Philosophy in 1808. In essence, he said that the atoms of a specific element are exactly alike and have the same weight; that the atoms of that specific element are different from the atoms of every other element; and that combinations of the elements are merely combinations of the atoms in simple or multiple units. At this time, Dalton did not know the exact number of atoms that could combine to form molecules. However, through experimentation, he knew the relative weights of the elements that formed compounds. From this information, he devel- oped the first table of atomic weights. Despite the fact that Dalton’s table was inaccurate, his concept was another step that advanced our knowledge of chemistry.

7 It became increasingly clear that atoms were the smallest, indivisible units of elementary

A Short History of Chemistry |

substances, but it was important to determine the atomic weights (mass) and sizes of differ- ent atoms. Since atoms were too minute to see with nineteenth-century microscopes, it was impossible to count and weigh them as we do with larger objects. Therefore, a different stan- dard was needed, and this is where the concept of “relative weights” became useful. Using this relative-weight concept makes it possible to determine the mass of an individual atom of one element relative to the mass of an atom of a different element. Using this system to ascertain atomic weight (mass), chemistry advanced rapidly. In a sense, all standards used for measure- ments of weight (pounds and kilograms), distance (miles and kilometers), volume (gallons and liters), temperature (Fahrenheit and Celsius), and so forth are arbitrarily related to a specific physical characteristic. Elements that are gases and consist of a single atom are called mona- tomic elements (Ar, Ne, Xe). Gas molecules that are composed of two atoms of an element

are called diatomic (O 2 ,H 2 , Cl 2 ), and three atoms of the same gaseous element can combine to form what are known as triatomic molecules (O 3 Ozone). The combinations of three or more atoms of different elements form polyatomic molecules (PO 4 ,H 2 SO 4 ). Thus, there are several distinct types of molecules. When atoms of elements combine to form molecules, the chemical and physical properties of the new molecules are different from the characteristics of the original elements. For instance, the compound NaCl is unlike the two elements that

make up this molecule. Na is sodium, a very reactive metal, and Cl 2 is chlorine, a reactive gas, whereas the molecule NaCl is well known as table salt. (There is further discussion of atomic structure later.) Based on what we have learned, let’s define elements and compounds.

Atom: The smallest unit of each of the more than 100 known elements and different basic types of matter that either exist in nature or are artificially made. All atoms that compose a specific element have the same nuclear charge and the same number of electrons and protons. Atoms of some elements may differ in mass when the number of neutrons in that atom’s nucleus is different (such atoms are called isotopes). This is discussed in the next chapter.

Compound: A molecular substance composed of atoms and formed by a chemical reaction of two or more elements. The molecules of the new compound have properties very different from the properties of the elements that formed the compound. An example is when two

diatomic molecules of hydrogen gas (2H 2 ) combine with one diatomic molecule of oxygen gas (O 2 ) to form two water molecules (2H 2 O), which have none of the properties of the original two substances. Allotrope: An allotrope is formed when an element or compound exists in more than one form. Carbon is an example of an element found in different forms (e.g., carbon black, graphite, and diamonds). Oxygen has three allotropes: monoatomic or nascent oxygen (O);

diatomic oxygen (O 2 ), the gas we breathe; and triatomic oxygen (O 3 ), which is known as ozone.

two

Atomic Structure

Early Ideas of Atomic Structure

As mentioned in the previous section titled “A Short History of Chemistry,” many scien- tists identified elements, determined their characteristics, similarities, and differences, and designed symbols for them. Using unique experiments, scientists devised ways to define the structure of atoms and determine atomic weights, sizes, and electrical charges as well as energy levels for atoms.

Many of these men and women recognized the existence of some order in the manner in which chemicals relate and react to each other. Although these scientists could not see the atoms themselves, they were aware that the structure of each element’s atoms has something to do with these characteristics. There were several attempts to organize the elements into a chart that reflected the particular nature of the atoms for these elements. Before the periodic table of the chemical elements was developed as we know it today, several relationships had to be established. (See the next section for more on the periodic table of the chemical elements.)

The concept of electrons had been known for many years, but determining how these negatively charged particles react required experimentation and analysis of data. In about 1897, Joseph John Thomson (1856–1940) sent streams of electrons through magnetic fields, which resulted in the dispersion or spreading of the electrons. Thomson’s experiments, and those of others, led him to speculate that the atom was a positively charged “core” and that negatively charged particles of energy surrounded and matched the positive charge of this core or nucleus. Further, when these electrons were excited or “stirred up” with strong light, electricity, or mag- netism, some of them were driven from the outer regions of the atom. This was one of the first experimental evidences for the structure of the atom. Many refinements of this concept were made by the work of several scientists, including the French chemist Marie Sklodwska Curie (1867–1943), the British physicist Baron Ernest Rutherford (1871–1937), the American physi- cist Robert Andrew Millikan (1868–1953), and others of many nationalities.

Two of Rutherford’s students conducted a classic experiment to determine the structure of an atom. They beamed alpha particles through a sheet of gold foil that was 1/50,000 of an inch thick. Since the thickness of this thin foil was only about two thousand atoms of gold, it

10 | The History and Use of Our Earth’s Chemical Elements should have allowed all the alpha particles to pass through and be detectable on the other side

of the foil. The experimenters recorded the pattern of alpha particles with a detecting instru- ment located behind the foil. Rutherford noticed that most particles went straight through the foil as if nothing interfered with them. However, unexpectedly, a few seemed to be diverted from the target. As it turned out, one out of every 10,000 alpha particles, as it passed through the foil, was deflected sideways, away from the center of the target located behind the foil, similar to a billiard ball glancing off another ball when struck. During the experiment, one out of every 20,000 particles bounced back from the foil toward the source of the alpha particles and thus did not pass through the gold foil at all. This indicated that some alpha particles were being deflected by something in the foil. After making some calculations, Rutherford concluded that this backward and side scattering of the alpha particles was evidence of a few collisions with something that had almost all its mass concentrated in a central, very small “nucleus.” He also determined that this tiny nucleus had a positive charge. Since the vast majority of the alpha particles passed straight through the foil to hit the center of the target used to detect the particles, the atoms of gold must be composed of mostly empty space. He stated, “It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper, and it came back to hit you.” * One way to understand the great distance between an atom’s nucleus and its surrounding electrons is to imagine a baseball on the floor of your living room as the nucleus of an atom. The electrons that surround the nucleus are about the size of peas that are orbiting in three dimensions at about a 10-mile radius from the baseball. This means that an atom with a 20-mile diameter would have a nucleus about the size of a baseball. The conclusion may be that there is more empty space than there is solid matter in matter and the universe, including in you.

Robert Andrew Millikan (1868–1953) continued the work of J. J. Thomson by conducting